1. Do now: Which species requires the least amount of energy to remove an electron from the outermost energy level? Na + F -1 O -2 Mg +2 Which substance will react with water, at room temperature and pressure to produce hydrogen? Mg NaH HN 3 NaOH NH 3

2. Bonding By: Thy Vuong-Schmick

3. Homework problems: Read Sections 8.1 – 8.5 Read Sections 8.6 – 8.8 Read Sections 9.1 – 9.5 Read Sections 9.6 – 9.8 Chapter 8 Chapter 9 1, 2, 35, 61, 2, 3, 4, 6, 78, 9, 10 18, 24, 2852, 54, 5514, 16, 20, 2250, 51, 54, 58 4, 30, 32, 3412, 57, 60, 6224, 26, 3060, 62, 63 36, 40, 4266, 68, 72, 80, 8632, 34, 36, 3866, 70, 71, 74 44, 46, 47, 50 40, 42, 44, 46, 48

4. Go over the test???

5. Bonding Activating strategy: Penny lab Crystal mini lab Hybridization Bond angle polarity

6. Do now: A compound contains only sulfur and oxygen is 50% sulfur by weight, what is the empirical formula for the compound? What is the mass of oxygen in 148 grams of calcium hydroxide?

7. The Science Behind the Fun Borax is an example of crystal solid with flat sides symmetrical shape because its molecules are arranged in a unique, repeating pattern Every crystal has a repeating pattern based on it's unique shape. They may be big or little, but they all have the same "shape“ Salt, sugar…are all examples of crystals How do the Borax crystals grow? Hot water holds more borax crystals than cold water heated water molecules move farther apart, making room for more of the borax crystals to dissolve. When no more of the solution can be dissolved, you have reached saturation. As this solution cools, the water molecules move closer together again. Now there's less room for the solution to hold onto as much of the dissolved borax. Crystals begin to form and build on one another as the water lets go of the excess and evaporates.

8. Bonding overview: Ions and electron configuration Ionic radii Ionization energy Higher ionization energies Electron affinity Ionic bonds and formation of ionic solids Alkali metals Alkaline earth metals Aluminum Halogens Noble gases Octet rule

9. Ions and electron configurations: Recall: Metals ______ ___ electrons and form ______________ Halogens and nometallic elements __________ electrons and form ________ Elements want to be noble gas Look at 1 st, 2 nd, halogen and oxygen group Aufbau principle: Metal in forming cation come from highest energy occupied orbital, while electrons that are accepted by nonmetal in forming an anion go into the lowest energy unoccupied orbital

11. Aufbau principle These rules are: 1.Electrons are placed in the lowest energetically available subshell. 2. An orbital can hold at most 2 electrons. 3.If two or more energetically equivalent orbitals are available (e.g., p, d etc.) then electrons should be spread out before they are paired up (Hund's rule).

12. Transition metals: Form cations by first losing valence shell “s” electrons and then losing “d” electrons Result: All remaining valence electrons in transition metal cations occupy “d” orbital Example: Fe +2 and Fe +3 Show configuration

13. Practice: Predict ground state electron configuration for each of the following ions: Ra +2 La +3 Ti +4 N -3 What doubly positive ion has the following ground state electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10

14. Radii: trend on the periodic table The cations are smaller than neutral atoms because: Principal quantum number of valence shell electron is smaller for the cations and because Z eff is larger

15. So go down a group: Size become ________ because:

16. Go across: Size become ____________ because:

17. Summary: Compare Na to Na +1 : The Na atom has 11 protons and 11 electrons Na +1 cation has 11 protons and 10 electrons. The smaller number of electrons in cation means that they shield one another to a lesser extent and therefore feel a stronger pull toward the nucleus

18. The opposite is true for anions

19. Practice problem: Which atom or ion in each of the following pairs would you expect to be larger? Explain. Oxygen atom and oxide ion Oxygen and sulfur atom Iron atom and iron (III) ion Hydrogen atom and hydride ion

20. Practice: Which of the following represents a K + ion, which is a K atom and which is a Cl -1 ion? R = 227 ppm R = 184 ppm R = 133 ppm

21. Revisit Ionization energy: Recall from last chapter: Absorption of light energy by atom leads to change in electron configuration Valence shell electron is promoted from lower energy to higher energy (larger quantum number n) If enough energy is absorbed: Electron can be removed completely from atom Leaving behind an ion Energy remove highest energy electron from an isolate neutral atom in “gaseous” state = Ionization energy

22. Ionization energy Always positive Energy required to remove electron from atom

23. Do now: What is the volume of 0.05 M HCl that is required to neutralize 50 mL of a 0.10 M Mg(OH) 2 ?

24. Talk about trend on the periodic table Ionization energy trend… Because…

25. Practice: unusual elements Compare between Be and B, ionization energy of these two elements? IE for B= 800 KJ IE for Be = 899 kJ WHY?????? 

26. Answer: -Draw out configuration for both -Because 2s electron is on average closer to the nucleus than 2p electron: -Held more tightly and harder to remove -IE of beryllium is larger than boron -Also because 2p electron of boron is shield by 2s electrons, feels a smaller Z eff and more easily removed than 2s electron of Be

27. Practice unusual elements: N and O Draw configuration to show removal of 1 electron Which has higher ionization energy? Why?

28. Diamagnetic atoms have only paired electrons, whereas paramagnetic atoms, which can be made magnetic, have at least one unpaired electron. Diamagnetic atoms repel magnetic fields. The unpaired electrons of paramagnetic atoms realign in response to external magnetic fields and are therefore attracted. Paramagnets do not retain magnetization in the absence of a magnetic field, because thermal energy randomizes electron spin orientations. Main difference between ferromagnets and paramagnets is that ferromagnets don't need outside magnetic field to become and stay magnetic.ferromagnets

29. Do now: Practice problem Which has larger 5 th ionization energy, Ge or As? Why? Which has larger 3 rd ionization energy, Be or N? Which has larger 4 th ionization energy, Ga or Ge?

30. Do now Three atoms have the following electron configurations: 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Which of those three has the largest E i1 ? Which has the smallest E i4 ?

31. Higher ionization energies: Two or three electrons can be removed from an atom Measured the amount of energy associated with each step: Larger amount of energy required for each ionization step Why?

32. Large jump in ionization energy Remember the chart on the test??? Large increase in ionization energies due to: Easy to remove an electron from a “partially filled” valence shell (Zeff is lower) Difficult to remove electron from “filled” valence shell (Zeff is higher) Example: Na, Mg, Al… Draw out the configuration and discuss the “jump”

33. Which has higher first ionization energy (E i1 ) N or O? Do not memorize the trend, because sometimes we have exceptions to rules This is an unusual case (because according to the trend, O should have higher ionization energy than N)

34. Practice problem: Three atoms have the following electron configurations: 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Which of those three has the largest E i1 ? Which has the smallest E i4 ?

35. Remember: Valence electrons are the one that lost during ionization Valence electrons also take part in bonding and chemical reactions Valence shell electron configuration of an atom controls the atom’s chemistry

36. Practice problems from power points H drive AP chem folder